-Chapter Five

I. History of the Periodic Table

• John Dalton(1776)-He suggested the first model of the atom with his atomic theory.
• John Dobereiner(1829)-He described the idea of a triad, which are groups of three elements with shared similar properties. Each triad has elements with the same number of valence electrons, so each triad contains elements with similar behavior.
• John Newlands(1827)-He arranged the existing elements by increasing atomic mass, and noticed that properties reemerged after every set of eight elements. He named this observatoin the law of octaves.
• Dmitri Mendeleev(1834)-He proposed the first true version of the periodic table. Considered to be the father of the periodic table, he was originally a Russian school teacher. Throughout his periodic table are empty spaces, which he reserved for elements he thought were yet to be discovered.
• Henry Moseley(1887)-He proposed the modern periodic table. He arranged the periodic table in order of increasing atomic number, as opposed to atomic mass like mendeleev had. He wrote the periodic law, which states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

  Today's Modern Table

Mendeleev's Original Table

2. Metals

• Metals are dactile, which means that they have the ability to be drawn into wires.
• Metals are Malleable as well, which means that they can be hammered into sheets.
• Metals have luster, which means that they reflect the light, giving off a shiny appearance.
• Metals are very accomplished conductors of electricity and heat, which means that heat and electricity run through them with ease.

3. Non-Metals

• Non-metals are brittle, which means that they are difficult to shape, and break easily when hammered.
• Non-metals have no luster, which means that they don't reflect light to emit a shiny apperance.
• Non metals are insulators, which means that they are unable to conduct electricity or heat.

4. Semi-Metals

• Semi metals have properties of both the Metals and Non-Metals accordingly.

Can YOU tell which elements are metals?

5. Period Trends

• There is a gradual decrease in atomic radii across the period from left to right
  • Ie: Lithium has an atomic radius of 152 compared to neon which has a radius of 71
• The increasing positive charge of the nucleus causes the trend of smaller atoms across a period.
• Ionization energies of the main-group elements increase across each period, because it is getting closer and closer to acquiring an electron and becoming a noble gas.
• Low ionizations energies lose electrons more easily and are therefore more reactive.
• Group 1 alkali metals are highly reactive because they have the lowest ionization energy
• As electrons add to atoms with increasing nuclear charge, electron affinities become more negative across each period.
• The halogens (Group 17) gain electrons most readily because they have large negative values of electron affinities. Because halogen atoms gain electrons so easily, Group 17 elements have high reactivities
• The metals at the left generally form cations and the nonmetals at the upper right generally form anions.
• Cationic radii decrease across a period because the electron cloud shrinks due to the increasing nuclear charge acting on the electrons in the same main energy level.
• Anionic radii decrease across each period for the elements in Groups 15-18.
• Electronegativities tend to increase across each period.
• Noble gases do not generally form compounds and therefore cannot be assigned electronegativities.
• The alkali and alkaline-earth metals are the least electronegative elements.
• Nitrogen, oxygen, and the halogens are the most electronegative elements.
• Electronegativities tend to either decrease down a group or remain about the same.
• The atomic radii of the main-group elements increase down a group because as electrons occupy sublevels in increasingly higher energy levels located farther from the nucleus, the sizes of the atoms increase.
  • Ie: Group 1- Hydrogen, 37, to Francium, 270
• Each successive electron removed from an ion feels an increasingly stronger effective nuclear charge.
• It is hardest to remove electrons from atoms in Group 18 elements.
• Generally, electrons add with great difficulty down a group. This is due to a slight increase in effective nuclear charge down a group, which increases electron affinities. There is also an increase in atomic radius down a group, which decreases electron affinities.
• Second electron affinities are all positive because it is always more difficult to add a second electron to an already negatively charged ion in the gas phase.
  • Ie: Chlorine has the configuration: [Ne]3s²3p^5
• There is a gradual increase of ionic radii down a group.

7. Ionization Energy/Valence Electrons

• Electrons can be removed from their atom if there is enough energy supplied.
• Process of removing electrons: A + energy à A+ + e- (A+ represents an ion of element A with a single positive charge.
• Valence electrons are most subject to the influence of nearby atoms or ions and are usually located in incompletely filled main-energy levels and are usually the electrons in the outermost s and p sublevels.
  • Ie: The electron lost from the 3s sublevel of Na to form Na+ is a valence electron.
• Each group has that number of valence electrons. Group 1 has one ve-, group 2 has 2 ve-. Groups 13-18 have a number of ve- equal to the group number minus 10.
• Linus Pauling devised a scale of numerical values reflecting the tendency of an atom to attract electrons.

8. Electronegativity/Atomic Radii

• High electron affinity=high electronegativity.
• The most electronegative element is fluorine with a value of four.The atomic radii of the d-block elements generally decrease across the periods.
• The radii increase because of repulsion among the electrons as the number of electrons in the d sublevel increases.

9. Ionization Energy

• Ionization energies of the d-block and f-block elements typically increase across the periods.
• The first ionization energies of the d-block elements tend to increase down each group because the electrons available for onization. in the outer sublevels are less shielded from the increasing nuclear charge by electrons in the incomplete (n-1)d sublevels.

10. Ion Formation and Ionic Radii

• Electrons in the highest occupied sublevel are always removed first out of all atoms of the d-block and f-block elements.
• The first electrons to be removed are those in the outermost s sublevels.
• The cations have smaller radii than the atoms do.
• The d-block elements all have electronegativities between 1.1 and 2.54
• The d-block elements follow the trend for electronegativity values to increase as radii decrease and vice versa.
• The f-block elements have similar electronegativities, which range from 1.1 to 1.5.